Your FREE AP Chemistry Flashcards 2026 – 200+ Cards

The SI unit for amount of substance; one mole = Avogadro's number of particles, $6.022\times10^{23}$.

$6.022\times10^{23}$ — the number of particles in one mole. Moles × this = number of particles.

Mass of one mole of a substance in g/mol, equal to the sum of the atomic masses in its formula.

Divide mass by molar mass: e.g. 88 g of $\ce{CO2}$ (44 g/mol) = 2 mol.

Multiply moles by Avogadro's number: 2 mol = $2\times6.022\times10^{23}=1.2\times10^{24}$ particles.

Separates isotopes by mass-to-charge ratio; peak positions give isotope masses and heights give relative abundances.

Atoms of the same element with different numbers of neutrons, hence different mass numbers.

The weighted average of an element's isotope masses by natural abundance; pulled toward the more abundant isotope.

Measures energy to remove electrons; each peak = a subshell, peak height ∝ number of electrons in it.

A filled s subshell (2 e⁻) vs a filled p subshell (6 e⁻) gives a 1-to-3 peak-height ratio.

The arrangement of electrons by subshell in order of increasing energy, e.g. Mg = $\text{1s}^2\text{2s}^2\text{2p}^6\text{3s}^2$.

Electrons fill the lowest-energy orbitals first before higher ones.

The net positive pull on a valence electron after shielding by core electrons; drives periodic trends.

Inner core electrons block valence electrons from the full nuclear charge because they lie between them and the nucleus.

s electrons penetrate closer to the nucleus than p electrons in the same shell, so they are lower in energy.

Increases down a group and to the left across a period (larger as you move away from the top-right).

Energy to remove an electron from a gaseous atom; increases across a period, decreases down a group.

How strongly an atom attracts shared bonding electrons; increases up and to the right — fluorine is highest.

Attraction increases with greater charge and decreases with greater distance — the basis of periodic trends.

A positively charged ion formed when an atom loses electrons; smaller than its parent atom.

A negatively charged ion formed when an atom gains electrons; larger than its parent atom.

The outermost-shell electrons that participate in bonding and determine chemical behavior.

Inner-shell electrons that shield valence electrons and do not participate in bonding.

The simplest whole-number ratio of atoms in a compound, found from percent composition by mass.

Electron transfer from a metal to a nonmetal forming oppositely charged ions; conducts when molten or dissolved.

A bond formed by sharing electron pairs between two nonmetals.

A 'sea' of delocalized electrons shared among metal cations; explains conductivity and malleability.

A diagram of a molecule's bonding and lone-pair electrons used to predict shape and reactivity.

Atoms tend to gain, lose, or share electrons to reach eight valence electrons (a noble-gas configuration).

An exception where the central atom has fewer than 8 electrons, e.g. boron in $\ce{BF3}$ (6 electrons).

An exception where a period-3+ central atom holds more than 8 electrons, e.g. $\ce{PCl5}$.

Electron domains around a central atom spread out to minimize repulsion, setting the molecular geometry.

A region of electron density — a bond (single, double, or triple) or a lone pair — counted in VSEPR.

2 electron domains, bond angle 180°, e.g. $\ce{CO2}$.

3 domains, 0 lone pairs, bond angle 120°, e.g. $\ce{BF3}$.

4 domains, 0 lone pairs, bond angle 109.5°, e.g. $\ce{CH4}$.

4 domains with 1 lone pair, angle ~107°, e.g. $\ce{NH3}$.

4 domains with 2 lone pairs, angle ~104.5°, e.g. $\ce{H2O}$.

5 bonding domains, angles 90° and 120°, e.g. $\ce{PCl5}$.

6 bonding domains, all 90° angles, e.g. $\ce{SF6}$.

5 domains with 1 lone pair (placed equatorial), e.g. $\ce{SF4}$.

6 domains with 2 lone pairs on opposite sides, e.g. $\ce{XeF4}$.

Head-on orbital overlap along the bond axis; every single bond is one sigma bond.

Side-by-side p-orbital overlap above and below the axis; double = 1 pi, triple = 2 pi.

The number of bonding electron pairs between two atoms; higher order = shorter, stronger bond.

Mixing of atomic orbitals: sp (linear), sp² (trigonal planar), sp³ (tetrahedral) to match geometry.

Unequal electron sharing from an electronegativity difference, creating a bond dipole.

Overall dipole from bond dipoles and geometry; symmetric $\ce{CO2}$ is nonpolar, bent $\ce{H2O}$ is polar.

Attractions between molecules that set boiling point, vapor pressure, and solubility.

Weak attractions from temporary, instantaneous dipoles; present in all molecules.

How easily an electron cloud distorts; larger clouds (more electrons) give stronger dispersion forces.

Attractions between the permanent dipoles of two polar molecules.

A strong IMF when H is bonded to F, O, or N; explains water's high boiling point.

Attraction between an ion and a polar molecule, e.g. $\ce{Na+}$ hydrated by water; strongest of the IMFs.

Ion-dipole > hydrogen bonding > dipole-dipole > London dispersion.

Stronger intermolecular forces require more energy to separate molecules, raising the boiling point.

Pressure of vapor in equilibrium with its liquid; lower when IMFs are stronger.

A liquid with weak IMFs, high vapor pressure, that evaporates readily.

Occurs when a liquid's vapor pressure equals the external atmospheric pressure.

A branched isomer has less surface contact, weaker dispersion forces, and a lower boiling point.

A state with particles in fixed positions, definite shape and volume; strongest particle interactions.

A state with definite volume but no fixed shape; particles touch but can move past one another.

A state with no fixed shape or volume; particles are far apart with negligible interactions.

$PV = nRT$ relates pressure, volume, moles, and temperature for an ideal gas.

Gas particles are in constant random motion; average kinetic energy is proportional to absolute temperature.

The pressure one gas in a mixture would exert alone; total pressure is the sum (Dalton's law).

Gases deviate from ideal behavior at high pressure and low temperature, where IMFs and molecular volume matter.

A homogeneous mixture of a solute dissolved in a solvent.

Polar solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes.

Concentration in moles of solute per liter of solution: $M = \text{mol}/\text{L}$.

Adding solvent lowers concentration while moles of solute stay constant: $M_1V_1 = M_2V_2$.

Two liquids that mix in all proportions, like water and ethanol.

Separating a mixture by differences in boiling point (vapor pressure).

Separating components by their differing attractions to a stationary vs a mobile phase.

Separating an insoluble solid from a liquid by passing the mixture through a porous barrier.

Absorbance is proportional to concentration and path length: $A = \varepsilon b c$.

Liquid rising in a narrow tube due to adhesion and cohesion (intermolecular forces).

Resistance of a liquid surface to increase its area, caused by cohesive intermolecular forces.

A liquid's resistance to flow; greater with stronger intermolecular forces.

A plot of pressure vs temperature showing the solid, liquid, and gas regions and their boundaries.

A substance going directly from solid to gas without passing through the liquid phase.

A solute that produces ions in solution and conducts electricity; strong electrolytes fully dissociate.

A process that rearranges atoms to form new substances, conserving mass and atoms.

An equation with equal numbers of each atom on both sides; coefficients give the mole ratio.

Using the mole ratios in a balanced equation to relate amounts of reactants and products.

The reactant that runs out first and caps the amount of product formed.

A reactant present in more than the stoichiometric amount; some is left over.

Actual yield ÷ theoretical yield × 100%, a measure of reaction efficiency.

Two or more substances combine into one product (A + B → AB).

One compound breaks down into two or more simpler substances.

One element replaces another in a compound; a redox reaction.

Ions exchange partners; often forms a precipitate, gas, or water.

A fuel reacts with oxygen, releasing heat; complete combustion gives $\ce{CO2}$ and $\ce{H2O}$.

Burning in limited oxygen produces carbon monoxide $\ce{CO}$ or soot instead of all $\ce{CO2}$.

Two solutions mix and form an insoluble solid (precipitate).

An ion that appears unchanged on both sides of an equation and is canceled in the net ionic equation.

An equation with all soluble strong electrolytes written as separate dissociated ions.

Shows only the species that actually react, after spectator ions are removed.

Loss of electrons; the oxidation number increases (OIL — Oxidation Is Loss).

Gain of electrons; the oxidation number decreases (RIG — Reduction Is Gain).

A bookkeeping charge assigned to an atom to track electron transfer in redox reactions.

The species that is reduced (gains electrons) and causes another to be oxidized.

The species that is oxidized (loses electrons) and causes another to be reduced.

Adding a known concentration of titrant until a reaction is complete to find an unknown amount.

How fast reactants are consumed or products formed; change in concentration over time.

Change in concentration of a species divided by the elapsed time interval.

$\text{rate}=k[A]^x[B]^y$; the orders x and y come from experiment, not coefficients.

The proportionality factor in a rate law; depends on temperature and activation energy, not concentration.

The exponent on a reactant in the rate law; the overall order is the sum of the exponents.

Rate is independent of concentration; a plot of [A] vs time is a straight line.

Rate ∝ [A]; a plot of ln[A] vs time is linear; units of k are $\text{s}^{-1}$.

Rate ∝ [A]²; a plot of $1/[A]$ vs time is linear.

Find order by changing one reactant: no change = 0, doubles = 1, quadruples (×4) = 2.

Time for a reactant's concentration to fall by half; constant for first-order reactions.

Reactions occur when particles collide with enough energy and correct orientation.

The minimum energy needed to react — the barrier between reactants and the transition state.

The high-energy 'activated complex' at the peak of the energy profile.

Relates the rate constant to temperature and activation energy; higher T or lower Eₐ speeds the reaction.

A substance that speeds a reaction via a lower-Eₐ pathway and is not consumed.

The step-by-step sequence of elementary steps by which a reaction occurs.

The slowest elementary step; it limits the overall reaction rate.

A species produced in one step and consumed in a later step; not in the overall equation.

Raising temperature increases the fraction of collisions with energy above Eₐ, speeding the reaction.

A catalyst speeds both directions equally; it does not shift equilibrium or change K.

The specific part of the universe being studied, such as the reacting chemicals.

Everything outside the system that can exchange energy with it.

Energy transferred because of a temperature difference; flows from warmer to cooler.

A measure of the average kinetic energy of particles, distinct from total heat.

A process that absorbs heat; $q$ and $\Delta H$ are positive (feels cold).

A process that releases heat; $q$ and $\Delta H$ are negative (feels hot).

Heat content at constant pressure; $\Delta H$ is the heat absorbed or released.

$\Delta H < 0$ = exothermic (products lower in energy); $\Delta H > 0$ = endothermic.

Energy to raise 1 g of a substance by 1 °C; water's is unusually high.

Energy to raise an entire object/sample by 1 °C (no per-gram basis).

$q = mc\Delta T$: heat = mass × specific heat × temperature change.

An insulated device that measures heat flow by tracking temperature change.

The heat absorbed or released when a solute dissolves in a solvent.

The $\Delta H$ of an overall reaction equals the sum of the step $\Delta H$ values (state function).

A property (like enthalpy) that depends only on the current state, not the path taken.

Energy to break a bond; $\Delta H$ ≈ bonds broken minus bonds formed.

$\Delta H_f^\circ$: heat change forming 1 mol of a compound from its elements in standard states.

The state where two objects in contact reach the same temperature and net heat flow stops.

A dynamic state where forward and reverse reaction rates are equal and concentrations stay constant.

Ratio of product to reactant concentrations at equilibrium, each raised to its coefficient.

A large K (≫1) favors products; a small K (≪1) favors reactants.

The same ratio as K but at any moment; comparing Q to K predicts the shift.

$Q<K$ shifts forward; $Q>K$ shifts reverse; $Q=K$ is at equilibrium.

Kc uses concentrations, Kp uses partial pressures; $K_p = K_c(RT)^{\Delta n}$.

Pure solids and liquids are left out of the K expression (their activity is 1).

A system at equilibrium shifts to partially offset an applied stress.

Shifts equilibrium toward products to consume the added reactant.

Shifts equilibrium toward products to replace what was removed.

Increasing pressure (lower volume) shifts toward the side with fewer moles of gas.

Only a temperature change actually changes the value of K; heating an exothermic reaction lowers K.

A catalyst reaches equilibrium faster but does not shift it or change K.

Equilibrium constant for a slightly soluble salt dissolving into its ions.

Moles of a salt that dissolve per liter; for a 1:1 salt, $K_{sp}=s^2$.

A precipitate forms when the ion product Q exceeds Ksp.

Adding an ion already in a solubility equilibrium decreases the salt's solubility.

A Reactants/Initial-Change-Equilibrium chart used to solve equilibrium problems.

A substance that produces $\ce{H+}$ (hydronium) ions in water.

A substance that produces $\ce{OH-}$ (hydroxide) ions in water.

A proton ($\ce{H+}$) donor.

A proton ($\ce{H+}$) acceptor.

Two species differing by one proton, e.g. $\ce{NH4+}$ and $\ce{NH3}$.

A strong acid fully ionizes; a weak acid only partially ionizes, setting up an equilibrium.

The stronger the acid, the weaker its conjugate base, which has little tendency to reaccept the proton.

A substance that can act as either an acid or a base, like water.

$\text{pH} = -\log[\ce{H+}]$; below 7 acidic, 7 neutral, above 7 basic.

$\text{pOH} = -\log[\ce{OH-}]$; at 25 °C, pH + pOH = 14.

The ion-product of water, $K_w = [\ce{H+}][\ce{OH-}] = 1.0\times10^{-14}$ at 25 °C.

Equals $-\log$ of its concentration, since a strong acid fully ionizes.

Acid dissociation constant; a larger Ka (smaller pKa) means a stronger acid.

$\text{p}K_a = -\log K_a$; a lower pKa means a stronger acid.

A weak acid and its conjugate base together; resists pH change on adding small amounts of acid or base.

$\text{pH} = \text{p}K_a + \log\frac{[\text{base}]}{[\text{acid}]}$.

When [base] = [acid] (1:1 ratio), the log term is 0, so pH = pKa.

The amount of acid or base a buffer can neutralize before pH changes significantly.

The point in a titration where moles of titrant exactly neutralize the analyte.

The region before the equivalence point where weak acid and conjugate base coexist and pH changes slowly.

Equivalence point pH is above 7 because the conjugate base is basic.

A substance that changes color over a pH range; choose one whose range spans the equivalence point.

An indicator colorless below ~8 and pink above ~10; good for weak-acid/strong-base titrations.

An acid and base react to form water and a salt; $\ce{H+ + OH- -> H2O}$.

A measure of the dispersal of matter and energy; the number of accessible microstates.

Forming more moles of gas increases disorder, so $\Delta S$ is positive.

A perfect crystalline solid at absolute zero (0 K) has zero entropy.

$\Delta G = \Delta H - T\Delta S$; the criterion for thermodynamic favorability.

A process with $\Delta G < 0$; favorable does not mean fast.

A favorable process driven mainly by a large negative $\Delta H$.

A favorable process driven mainly by a large positive $\Delta S$.

Where favorability flips; set $\Delta G = 0$, so $T = \Delta H/\Delta S$.

$\Delta G^\circ = -RT\ln K$; a negative $\Delta G^\circ$ corresponds to $K>1$.

Uses a spontaneous redox reaction to produce electricity.

The electrode where oxidation occurs (loss of electrons).

The electrode where reduction occurs (gain of electrons).

The reference half-cell defined as exactly $0$ V.

$E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$; positive means spontaneous.

$\Delta G^\circ = -nFE^\circ$; a larger positive $E^\circ$ gives a more negative $\Delta G^\circ$.

Uses an external power source to drive a non-spontaneous redox reaction.